Chapters

Everything NCERT Class 11 Unit 3 (Classification of Elements and Periodicity in Properties) says that can be asked, plus what JEE Main adds beyond the book (marked Beyond NCERT). The short note stays for the last week; this is for a proper revision pass.

Formula sheet

The chapter has few formulas; its rules and trends fit on one screen. Details and examples are in the sections below.

Quantity	Rule	Watch
Period	= n of the valence shell	Elements in a period = 2 × orbitals being filled
Group, s-block	= number of ns electrons	He is s² but sits in group 18
Group, d-block	= (n − 1)d + ns electrons	Pd 4d¹⁰5s⁰ is group 10
Group, p-block	= 10 + ns + np electrons
Valence (main group)	= valence electrons, or 8 minus them	Group 15: 3 or 5
Covalent radius	½ × single-bond length in X₂	Cl₂ 198 pm → 99 pm
Metallic radius	½ × distance between metal cores	Cu 256 pm → 128 pm
Ionisation enthalpy	X(g) → X⁺(g) + e⁻; always positive	IE₁ < IE₂ < IE₃ …
Electron gain enthalpy	X(g) + e⁻ → X⁻(g); sign varies	ΔegH = −Aₑ − (5/2)RT
IUPAC name for Z > 100	digit roots + "ium"	0 nil, 1 un, 2 bi … 9 enn

Property	Across a period →	Down a group ↓
Atomic and ionic radius	Decreases	Increases
Ionisation enthalpy	Increases	Decreases
Electron gain enthalpy	More negative	Less negative
Electronegativity	Increases	Decreases
Metallic character	Decreases	Increases
Basic character of oxides	Decreases	Increases

0 · What JEE Main asks from this chapter

131 questions in JEE Main 2019–2026, about 1 per paper; nearly all multiple-choice (126), most of them orders and statements. Grouped by type (counts are approximate):

Type	Share	Typical question
Ionisation enthalpy: orders, anomalies, successive IEs	~24%	Increasing IE₁ of Mg, Al, S, P, Si (2021)
Position and naming: block, group, period, IUPAC names, history	~21%	IUPAC symbol for Z = 119 (2019)
Atomic and ionic radii, isoelectronic species	~19%	Order of O²⁻, F⁻, Na⁺, Mg²⁺ (2020)
Oxides, metallic character, reactivity	~15%	Acidic, basic and amphoteric oxides from a list (2020)
Electron gain enthalpy	~13%	Order for Cl, F, Te, Po (2022)
Electronegativity	~8%	Correct Pauling comparison: Ga < Ge (2019)

Almost every question is an order of three to five elements with one exception planted in it. Learn the exceptions in sections 7–10 cold.

1 · Genesis of the periodic table

In 1800, 31 elements were known; by 1865, 63. NCERT's opening says 114 are known; its Table 3.5 names all 118. Classifying them organises known facts and predicts new ones.

1. Döbereiner's triads (1829): groups of three similar elements; the middle one's atomic weight is about the mean of the other two. Li 7, Na 23, K 39; Ca 40, Sr 88, Ba 137; Cl 35.5, Br 80, I 127. Worked for only a few elements; dismissed as coincidence.
2. de Chancourtois (French geologist, 1862): elements by atomic weight on a cylinder, showing periodic recurrence. Ignored.
3. Newlands' law of octaves (1865): elements by atomic weight; every eighth resembles the first, like octaves in music. True only up to calcium. Davy Medal, 1887.
4. Lothar Meyer (1869, table by 1868): plotted atomic volume, melting point and boiling point against atomic weight and found a repeating pattern whose length changes. Published after Mendeleev.
5. Mendeleev (1869): first to publish the periodic law: the properties of the elements are a periodic function of their atomic weights.

(2025: "Newlands arranged elements by atomic number" is false: by atomic weight. "Meyer plotted against atomic number" is also false: atomic weight.)

Mendeleev's table:

• Rows (series) and columns (groups) by increasing atomic weight; similar elements in one group. He used a wide range of properties, especially the formulas of compounds.
• Where weights disagreed with properties, he followed properties: iodine (lighter) went after tellurium, into group VII with F, Cl, Br.
• He left gaps for undiscovered elements and predicted their properties: eka-aluminium (gallium) and eka-silicon (germanium).

Property	Eka-Al → Ga (predicted → found)	Eka-Si → Ge (predicted → found)
Atomic weight	68 → 70	72 → 72.6
Density (g cm⁻³)	5.9 → 5.94	5.5 → 5.36
Melting point (K)	low → 302.93	high → 1231
Oxide	E₂O₃ → Ga₂O₃	EO₂ → GeO₂
Chloride	ECl₃ → GaCl₃	ECl₄ → GeCl₄

2 · The modern periodic law and the long form

Moseley (1913) studied the characteristic X-ray spectra of elements. A plot of √ν against atomic number Z is a straight line; √ν against atomic mass is not. So atomic number is the more fundamental property.

(The equation is Beyond NCERT; the straight line of √ν vs Z is NCERT. Asked 2023 and 2025: ν itself vs Z is not a straight line.)

Modern periodic law: the physical and chemical properties of the elements are periodic functions of their atomic numbers. Periodicity is really the periodic repetition of electronic configurations.

• 94 elements occur naturally (Np and Pu, like Ac and Pa, occur in pitchblende, an ore of uranium).
• Long form: horizontal rows are periods (7), vertical columns are groups (18, IUPAC numbering 1–18, replacing IA … VIIA, VIII, IB … VIIB and 0).
• Period lengths: 2, 8, 8, 18, 18, 32, 32 (the 7th now complete at 118).
• The 14 lanthanoids and 14 actinoids sit in separate rows at the bottom, to keep the table's structure and keep similar elements in one column.

3 · Naming elements with Z > 100

Discoverers used to name new elements; disputes followed (element 104: Rutherfordium for the Americans, Kurchatovium for the Soviets). IUPAC's rule: until a discovery is proved and named, use a systematic name from the digits.

Roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn. Join them in the order of the digits and add -ium; the symbol is the roots' first letters (three letters). Bi and tri lose their final i before -ium (Unnilbium, Unniltrium).

Example (NCERT Problem 3.1) · Z = 120: un-bi-nil-ium → Unbinilium, Ubn. Z = 119: Ununennium, Uue (2019). Unnilennium = 109 (2020). Unnilunium = 101 (2020).

Official names (Table 3.5): 101 Mendelevium Md, 102 Nobelium No, 103 Lawrencium Lr, 104 Rutherfordium Rf, 105 Dubnium Db, 106 Seaborgium Sg, 107 Bohrium Bh, 108 Hassium Hs, 109 Meitnerium Mt, 110 Darmstadtium Ds, 111 Roentgenium Rg, 112 Copernicium Cn, 113 Nihonium Nh, 114 Flerovium Fl, 115 Moscovium Mc, 116 Livermorium Lv, 117 Tennessine Ts, 118 Oganesson Og.

Example (2025) · Families of the newest elements: pnictogen (group 15) Mc, chalcogen (16) Lv, halogen (17) Ts, noble gas (18) Og.

Example (2024) · Unununnium = 111 → [Rn]5f¹⁴6d¹⁰7s¹ → group 11. An element [Rn]5f¹⁴6d¹7s² is Z = 103, Unniltrium (2022).

4 · Electronic configuration and position

An element's place reflects the quantum numbers of the last orbital filled.

Periods: the period number is n of the valence shell. Each period holds twice the number of orbitals being filled.

Period	Orbitals filled	Elements
1	1s	2 (H, He)
2	2s, 2p	8 (Li to Ne)
3	3s, 3p	8 (Na to Ar)
4	4s, 3d, 4p	18 (K to Kr; 3d series Sc 21 to Zn 30)
5	5s, 4d, 5p	18 (Rb to Xe; 4d series from Y 39)
6	6s, 4f, 5d, 6p	32 (Cs to Rn; lanthanoids Ce 58 to Lu 71)
7	7s, 5f, 6d, 7p	32 (Fr to Og; actinoids Th 90 to Lr 103)

Example (NCERT Problem 3.2) · Why 18 in period 5: it fills 5s, 4d, 5p (in that order of energy) = 1 + 5 + 3 = 9 orbitals → 18 electrons. Period 6 (exercise 3.4) fills 6s, 4f, 5d, 6p = 16 orbitals → 32.

Groups: elements in a group have the same number and arrangement of outer electrons, so similar properties (group 1: [He]2s¹, [Ne]3s¹ … [Rn]7s¹).

Finding position from Z · Beyond NCERT method: find the noble gas just below Z; the remainder fills ns, then (n − 2)f, (n − 1)d, np.

Example (NCERT Problem 3.3, 2019) · Z = 117: [Rn]5f¹⁴6d¹⁰7s²7p⁵ → group 17. Z = 120: [Og]8s² → group 2, an alkaline earth metal.

Example (2019) · Z = 71 (Lu, [Xe]4f¹⁴5d¹6s²): the 71st electron enters 5d. Z = 101 is an actinoid, Z = 104 is group 4 (2020).

5 · The four blocks; metals, non-metals, metalloids

s-block (groups 1, 2; ns¹ and ns²):

• Reactive metals with low ionisation enthalpies; form M⁺ (alkali metals) or M²⁺ (alkaline earth metals).
• Metallic character and reactivity increase down the group.
• Never found pure in nature (too reactive; 2024 statement).
• Compounds are mainly ionic, except those of Li and Be.

p-block (groups 13–18; ns²np¹ to ns²np⁶). With the s-block these are the representative or main group elements.

• Each period ends in a noble gas, ns²np⁶, closed shell, very low reactivity.
• Before them: halogens (17) and chalcogens (16), with highly negative electron gain enthalpies; they gain one or two electrons to reach the noble gas configuration.
• Non-metallic character rises left to right; metallic character rises down a group.

d-block (groups 3–12; (n − 1)d¹⁻¹⁰ns⁰⁻², Pd 4d¹⁰5s⁰): the transition elements.

• All metals; mostly coloured ions, variable oxidation states, paramagnetism, used as catalysts.
• Zn, Cd, Hg ((n − 1)d¹⁰ns²) lack most of these properties.
• They bridge the reactive s-block metals and the less active groups 13 and 14, hence "transition".

f-block (inner transition): lanthanoids Ce (58) to Lu (71), actinoids Th (90) to Lr (103); (n − 2)f¹⁻¹⁴(n − 1)d⁰⁻¹ns².

• All metals; properties within each series are very similar.
• Early actinoids have more oxidation states, so their chemistry is more complicated than the lanthanoids'.
• Actinoids are radioactive; many are made only in nanogram amounts. Elements after uranium are transuranium elements.

Two exceptions in placing:

• Helium is s-block (1s²) but sits in group 18: its shell is full and it behaves as a noble gas.
• Hydrogen has one s electron (like group 1) but can gain one electron (like group 17), so it is placed separately at the top.

Metals, non-metals, metalloids:

• Metals are over 78% of elements, on the left. Solids at room temperature except mercury; gallium (303 K) and caesium (302 K) melt just above room temperature. High m.p. and b.p., good conductors, malleable, ductile.
• Non-metals (fewer than twenty) are at the top right. Solids or gases with low m.p. and b.p. (boron and carbon are exceptions), poor conductors, brittle.
• Metallic character falls across a period and rises down a group. The change is gradual: the elements along the diagonal zig-zag line, Si, Ge, As, Sb, Te, are metalloids (semi-metals).

Example (NCERT Problem 3.4, asked 2022) · Increasing metallic character: P < Si < Be < Mg < Na.

Example (2021) · Z = 33 (As) metalloid, 53 (I) non-metal, 83 (Bi) metal. Te is the metalloid among Sc, Pb, Bi, Te (2022).

6 · Atomic radius

An atom's size (about 1.2 Å) can't be measured directly: it is tiny and its electron cloud has no sharp edge. It is estimated from distances between atoms that are bonded.

• Covalent radius (non-metals): half the single-bond distance. Cl₂ is 198 pm, so Cl is 99 pm.
• Metallic radius (metals): half the distance between neighbouring metal cores. Cu–Cu is 256 pm, so Cu is 128 pm.
• NCERT calls both "atomic radius". Measured by X-ray or spectroscopic methods.
• Noble gases are left out: being monoatomic, they have only van der Waals radii, which are much larger. Compare them only with other van der Waals radii.

(2022: "in Cl₂ the covalent radius is double the atomic radius" is false: the bond length is double.)

Trends:

• Across a period: decreases. Same valence shell, but Zeff rises, so electrons are pulled in.
• Down a group: increases. n rises; inner filled shells shield the valence electrons.

Atomic radii in pm (NCERT Table 3.6):

• Period 2: Li 152, Be 111, B 88, C 77, N 74, O 66, F 64.
• Period 3: Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99.
• Group 1: Li 152, Na 186, K 231, Rb 244, Cs 262.
• Group 17: F 64, Cl 99, Br 114, I 133, At 140.

Smaller-than-expected radii · Beyond NCERT:

• Ga (135 pm) is slightly smaller than Al (143 pm): the ten 3d electrons shield poorly.
• Zr ≈ Hf, Nb ≈ Ta, Mo ≈ W: the lanthanoid contraction (4f shields poorly).

Example (2019, 2020) · C < S < Al < Cs. F < O < C < Cl < Br.

7 · Ionic radius and isoelectronic species

Ionic radii come from cation–anion distances in ionic crystals; they follow the same trends as atomic radii.

• Cation < parent atom: fewer electrons, same nuclear charge. Na 186 pm → Na⁺ 95 pm (NCERT's text; its own table gives 102 pm).
• Anion > parent atom: the extra electron raises repulsion and lowers Zeff. F 64 pm → F⁻ 136 pm.

(2024 and 2026: "ions are always smaller than their atoms" and "atomic radius always exceeds ionic radius" are both false: anions are larger.)

Isoelectronic species have the same number of electrons. Their size is decided by nuclear charge (2019, 2024): more protons pull the same electrons closer.

• Greater positive charge → smaller; greater negative charge → larger.

Values worth knowing (pm): N³⁻ 171, O²⁻ 140, F⁻ 133, Na⁺ 102, Mg²⁺ 72, Al³⁺ 54.

Example (NCERT Problem 3.5) · Mg, Mg²⁺, Al, Al³⁺: largest Mg, smallest Al³⁺.

Example (2023) · Isoelectronic set: K⁺, Cl⁻, Ca²⁺, Sc³⁺ (18 each). From O²⁻, F⁻, Al, Mg²⁺, Na⁺, O⁺, Mg, Al³⁺, F, the 10-electron ones are O²⁻, F⁻, Mg²⁺, Na⁺, Al³⁺ = 5.

Example (2021) · Na⁺ 1.02 Å, so Mg²⁺ and Al³⁺ are 0.72 and 0.54 Å. N³⁻ is bigger than O²⁻ (1.40 Å) and F⁻ (1.33 Å), and far bigger than covalent N (0.74 Å).

8 · Ionisation enthalpy

Ionisation enthalpy (ΔᵢH, kJ mol⁻¹): energy to remove an electron from an isolated gaseous atom in its ground state. ("Isolated gaseous" and "ground state" make values comparable.) Unqualified, it means the first.

• Always positive (endothermic).
• IE₂ > IE₁: it is harder to remove an electron from a positive ion. IE₃ > IE₂, and so on.
• On the plot of IE₁ against Z, maxima at noble gases (closed shells), minima at alkali metals (hence their reactivity).

Two factors: attraction of the nucleus, and repulsion between electrons. Inner electrons shield (screen) the valence electron, so it feels Zeff, less than Z. Shielding works best when inner shells are full, as in the alkali metals.

• Across a period: increases. Electrons go into the same shell; the shielding added doesn't keep up with the rising nuclear charge.
• Down a group: decreases. The valence electron is farther out and more shielded; shielding outweighs the larger nuclear charge.

The two dips (NCERT explains both):

• B < Be: B loses a 2p electron, Be a 2s. A 2s electron penetrates closer to the nucleus, so the 2p electron of B is more shielded and easier to remove.
• O < N: N's three 2p electrons are in separate orbitals (Hund). In O, two of the four 2p electrons share an orbital; the extra electron–electron repulsion makes one easier to remove.

So period 2 runs: Li < B < Be < C < O < N < F < Ne. Period 3 does the same: Na < Al < Mg < Si < S < P < Cl < Ar (496, 577, 737, 786, 1000, 1012, 1251, 1521).

Example (NCERT Problem 3.6, asked 2020, 2022) · Na 496, Mg 737, Si 786: Al is nearer 575 than 760, because its 3p electron is shielded by the 3s electrons. (Actual 577.)

Example (2021, 2026) · Increasing IE₁: Al < Mg < Si < S < P; third period trend Al < Si < S < P < Cl.

Successive ionisation enthalpies. A huge jump marks the end of the valence electrons: the next electron comes from a noble-gas core.

Example (2020, 2025) · 800, 2427, 3658, 25 024, 32 824 kJ mol⁻¹: the jump comes after the third → 3 valence electrons → group 13 (boron).

Example (2020) · IE₁ 496, IE₂ 4560: one valence electron, an alkali metal (Na). 1 mol MOH needs 1 mol HCl or 0.5 mol H₂SO₄.

Na vs Mg (NCERT exercise 3.17; 2026): IE₁(Na) < IE₁(Mg), but IE₂(Na) > IE₂(Mg): Na⁺ is [Ne], so its second electron comes from a closed shell; Mg⁺ still has a 3s electron. For the same reason K has the greatest IE₂ − IE₁ gap among Ba, K, Ca, Sc (2019), and among Ca, Al, Fe, B it is Ca whose IE₃ is very high (2026).

Second ionisation enthalpies · Beyond NCERT (2026): compare the cations' configurations. C⁺ 2p¹, N⁺ 2p², O⁺ 2p³ (half-filled), F⁺ 2p⁴. Order C < N < F < O.

Group 13 and 14 irregularities · Beyond NCERT (NCERT exercise 3.19): B 801, Al 577, Ga 579, In 558, Tl 589 kJ mol⁻¹. Poorly shielding d (and f) electrons make Ga and Tl hold their electrons harder. In group 14, Pb > Sn.

Isotopes of one element have the same IE (same electrons, same nuclear charge; NCERT exercise 3.25).

9 · Electron gain enthalpy

Electron gain enthalpy (ΔegH): the enthalpy change when an electron is added to an isolated gaseous atom.

• It can be negative (energy released) or positive (energy needed).
• Halogens: very negative; one electron gives a noble gas configuration.
• Noble gases: large positive; the electron must enter the next shell.
• Most negative at the upper right, just before the noble gases.
• Less regular than ionisation enthalpy.

Trends:

• Across a period: generally more negative (larger Zeff, smaller atom).
• Down a group: generally less negative (bigger atom).
• Exception: O and F are less negative than S and Cl. The added electron enters the small n = 2 shell and is repelled by the electrons already there; in the larger n = 3 shell repulsion is much less.

ΔegH in kJ mol⁻¹ (NCERT Table 3.7):

• Group 1: H −73, Li −60, Na −53, K −48, Rb −47, Cs −46.

• Group 16: O −141, S −200, Se −195, Te −190, Po −174.

• Group 17: F −328, Cl −349, Br −325, I −295, At −270.

• Group 18: He +48, Ne +116, Ar +96, Kr +96, Xe +77, Rn +68.

• Most negative of all: Cl. Halogen order by magnitude: Cl > F > Br > I. Group 16: S > Se > Te > Po > O.

• Nearly equal pairs (2022): Rb and Cs; Ar and Kr.

Example (NCERT Problem 3.7) · Among P, S, Cl, F: most negative Cl, least negative P.

Example (2024) · Increasing negativity: Ar (positive) < S < Br < F, so ΔegH order Ar > S > Br > F.

Second electron gain enthalpy (NCERT exercise 3.21; 2019): positive. O⁻ + e⁻ → O²⁻ forces an electron onto a negative ion. So O²⁻ forms in oxides only because lattice energy pays for it.

Positive (endothermic) electron gain · Beyond NCERT: Be and Mg (full s), N (half-filled p; about zero), noble gases, and every second electron gain. The gain to H is exothermic (2020).

Electron affinity (NCERT footnote): many books use Aₑ, the negative of ΔegH, so a released energy gives a positive Aₑ. Strictly it is defined at absolute zero:

(2024's "electron affinity is negative for Be → Be⁻, N → N⁻, O → O²⁻" uses this sign: those three need energy.)

10 · Electronegativity

Electronegativity: a qualitative measure of an atom's ability, in a compound, to attract the shared electrons to itself.

• Not measurable, unlike IE and ΔegH. Scales: Pauling (most used), Mulliken–Jaffé, Allred–Rochow.
• Pauling (1922) set F = 4.0, the most electronegative.
• An element's value isn't constant: it depends on the atom it is bonded to (2023, 2025: "C is 2.5 in both CH₄ and CCl₄" is false).
• Across a period: increases; down a group: decreases, like IE, and opposite to atomic radius.
• Higher electronegativity goes with non-metallic character; lower with metallic.

Pauling values (NCERT Table 3.8):

• Period 2: Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0 (steps of 0.5).
• Period 3: Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3.0.
• Group 1: Li 1.0, Na 0.9, K 0.8, Rb 0.8, Cs 0.7.
• Group 17: F 4.0, Cl 3.0, Br 2.8, I 2.5, At 2.2.

Example (2019) · Al's electronegativity (1.5) equals Be's (diagonal relationship). Ga < Ge is the correct comparison; Te > Se, Si < Al and P > S are all wrong.

Example (2026) · Group 15: N 3.0, P 2.1, As 2.0. N − P = 0.9 > P − As = 0.1 → X = N, Y = As.

Which properties fall down a group? Electronegativity, IE and |ΔegH| (generally) fall; atomic radius and metallic character rise (2019). Across a period, only atomic radius (and metallic character) runs the other way (2020).

11 · Valence, oxidation state, anomalous second period

Valence of a representative element = number of valence electrons, or 8 minus it.

Group	Valence electrons	Valence
1	1	1
2	2	2
13	3	3
14	4	4
15	5	3, 5
16	6	2, 6
17	7	1, 7
18	8	0, 8

NCERT's Table 3.9 shows the pattern in formulas: hydrides LiH, CaH₂, B₂H₆, CH₄, NH₃, H₂O, HF; oxides Li₂O, MgO, B₂O₃, CO₂, N₂O₅ (and N₂O₃), SO₃, Cl₂O₇.

Example (NCERT Problem 3.8) · Si (valence 4) + Br (1) → SiBr₄; Al (3) + S (2) → Al₂S₃.

Oxidation state: the charge an atom takes on the basis of electronegativity. In OF₂, F (more electronegative) is −1, so O is +2; in Na₂O, O is −2, Na +1 (2026 statement).

Example (NCERT Problem 3.9) · In [AlCl(H₂O)₅]²⁺ the oxidation state of Al is +3 but its covalency is 6: they differ.

Anomalous first members. Li, Be and B to F differ from the rest of their groups. Li and Be form compounds with marked covalent character, while their heavier relatives are mainly ionic. Reasons:

1. Small size, large charge/radius ratio, high electronegativity.
2. Only four valence orbitals (2s, 2p), so maximum covalency 4: B forms BF₄⁻, but Al forms AlF₆³⁻.
3. Strong pπ–pπ multiple bonds to themselves (C=C, C≡C, N=N, N≡N) and to other period 2 elements (C=O, C=N, C≡N, N=O).

Example (2024) · Of B, C, N, S, O, F, P, Al, Si, three can't reach their group valency: N, O and F (no d orbitals, covalency ≤ 4; F and O are the most electronegative).

Diagonal relationship: a period 2 element resembles the period 3 element one group to its right: Li–Mg, Be–Al, B–Si. Li (152 pm) and Mg (160 pm) are close in size; Li⁺ 76 pm vs Mg²⁺ 72 pm. (Li–Na in 2021 is the odd pair out: they are a group relationship.)

Example (2021) · E is 4s²4p¹ (Ga). The p-block period-5 element diagonal to it (one group right, one period down) is Sn: [Kr]4d¹⁰5s²5p².

12 · Periodic trends and chemical reactivity

Every property follows from the electronic configuration.

• Across a period, radius falls, IE rises and ΔegH becomes more negative. So reactivity is highest at the two ends and lowest in the middle:
  • left end (alkali metals): by losing an electron (low IE), the reducing and metallic side;
  • right end (halogens): by gaining one (most negative ΔegH), the oxidising and non-metallic side.
• (2024: "reactivity rises steadily from group 1 to 18" is false.)
• Group 1 reactivity: Li < Na < K < Rb < Cs (easier loss). Group 17: F > Cl > Br > I (exercise 3.28).
• Transition metals: IEs between the s- and p-blocks, so less electropositive than groups 1 and 2; down a transition group the trend reverses.

Oxides (reaction with oxygen shows reactivity best):

Kind	Where	Examples
Basic	Left end	Na₂O, K₂O, Li₂O, CaO, BaO, MgO
Amphoteric	Middle	Al₂O₃, As₂O₃ (NCERT); BeO, ZnO, SnO₂, PbO₂, Ga₂O₃, Cr₂O₃
Acidic	Right end	Cl₂O₇, SO₃, P₄O₁₀, N₂O₅, N₂O₃, CO₂, SiO₂, B₂O₃, Mn₂O₇
Neutral	Middle	CO, NO, N₂O

• Amphoteric oxides act as acids with bases and as bases with acids; neutral oxides do neither.
• Basic character: K₂O > Na₂O > MgO > Al₂O₃ (2026).
• A metal's lowest oxidation state gives the most basic oxide (V₂O₃ basic); its highest gives an acidic one (Mn₂O₇, CrO₃) · Beyond NCERT.

(NCERT Problem 3.10: a strong base and a strong acid.)

Example (2020) · Acidic, basic, amphoteric in that order: N₂O₃, Li₂O, Al₂O₃. Period 3 elements whose oxides are basic, amphoteric, acidic (X, Y, Z): X < Y < Z in atomic number.

13 · Traps

Trap	What goes wrong	Do this
Most negative ΔegH	Answering F	Cl (−349) beats F (−328); S beats O
IE₁ in a period	Following the plain rise	Be > B, N > O; Mg > Al, P > S
Isoelectronic size	Thinking equal electrons, equal size	More protons, smaller; decided by nuclear charge
"Ions are smaller"	Applying it to anions	Anions are larger than their atoms
Second ΔegH	Thinking O⁻ + e⁻ releases energy	Positive: the ion repels the electron
Electron affinity sign	Mixing Aₑ and ΔegH	Aₑ = −ΔegH (roughly); read which one
Ga vs Al	Expecting Ga larger and easier to ionise	Ga 135 < Al 143 pm; IE Ga 579 > Al 577
Newlands, Meyer	Saying atomic number	Both used atomic weight
Moseley	ν vs Z as the straight line	√ν vs Z is straight
IUPAC symbols	Two letters, or roots out of order	One letter per digit, in order: 119 Uue, 120 Ubn

14 · Numbers worth knowing

• IE₁ (kJ mol⁻¹): H 1312, He 2372; Li 520, Be 899, B 801, C 1086, N 1402, O 1314, F 1681, Ne 2080; Na 496, Mg 737, Al 577, Si 786; K 419.
• Li IE₂ = 7297 (a 14-fold jump); Mg IE₂ = 1450, Na IE₂ = 4560.
• ΔegH: F −328, Cl −349, Br −325, I −295; O −141, S −200.
• Pauling: F 4.0, O 3.5, N 3.0 = Cl 3.0, Br 2.8, C 2.5 = S 2.5 = I 2.5, H 2.1 · (H value Beyond NCERT).
• Radii (pm): Na 186, Na⁺ 102; F 64, F⁻ 133–136; Cl 99; Mg²⁺ 72; Al³⁺ 54; O²⁻ 140.

15 · For the boards only

JEE Main doesn't ask this, but the CBSE paper can.

• Seaborg (Nobel 1951) discovered plutonium (1940) and the transuranium elements 94–102, and moved the actinoids below the lanthanoids; element 106, seaborgium, honours him.
• Mendeleev wrote Principles of Chemistry; element 101 honours him. The structure of the atom was unknown when he made his table (a 2021 statement).